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Acids and Bases

Chapter 15

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Properties of Acids• Sour taste• Change color of vegetable dyes• React with “active” metals

– Like Al, Zn, Fe, but not Cu, Ag or Au

Zn + 2 HCl ZnCl2 + H2

– Corrosive

• React with carbonates, producing CO2

– Marble, baking soda, chalk

CaCO3 + 2 HCl CaCl2 + CO2 + H2O

• React with bases to form ionic salts – And often water

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Properties of Bases

• Also Known As Alkalis• Taste bitter• Feel slippery• Change color of vegetable dyes

– Different color than acid

– Litmus = blue

• React with acids to form ionic salts– And often water

– Neutralization

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Arrhenius Theory• Acids ionize in water to H+1 ions and anions• Bases ionize in water to OH-1 ions and cations• Neutralization reaction involves H+1 combining

with OH-1 to make water• H+ ions are protons• Definition only good in water solution• Definition does not explain why ammonia

solutions turn litmus blue– Basic without OH- ions

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Brønsted-Lowery Theory• H+1 transfer reaction

– Since H+1 is a proton, also known as proton transfer reactions• Acid is H+ donor; Base is H+ acceptor

– Base must contain an unshared pair of electrons• In the reaction, a proton from the acid molecule is

transferred to the base molecule– H forms a bond to lone pair electrons on the base molecule– We consider only 1 H transferred in each reaction

• Products are called the Conjugate Acid and Conjugate Base– After reaction, the original acid is the conjugate base and the

original base is changed to what is now called the conjugate acid

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Brønsted-Lowery TheoryH-A + :B A-1 + H-B+1

A-1 is the conjugate base, H-B+1 is the conjugate acid• Conjugate Acid-Base Pair is either the original

acid and its conjugate base or the original base and its conjugate acid– H-A and A-1 are a conjugate acid-base pair– :B and H-B+1 are a conjugate acid-base pair

• The conjugate base is always more negative than the original acid; and the conjugate acid is always more positive than the original base

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Example #1

• Determine what species you will get if you remove 1 H+1 from the acid

– The Conjugate Base will have one more negative charge than the original acid

H3PO4 H+1 + H2PO4-1

Write the conjugate base for the acid H3PO4

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Brønsted-Lowery Theory

• In this theory, instead of the acid, HA, dissociating into H+1(aq) and A-1(aq); The acid donates its H to a water molecule

HA + H2O A-1 + H3O+1

A-1 is the conjugate base, H3O+1 is the conjugate acid• H3O+1 is called hydronium ion• In this theory, substances that do not have OH-1 ions

can act as a base if they can accept a H+1 from water

H2O + :B OH-1 + H-B+1

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Strength of Acids & Bases• The stronger the acid, the more willing it is to donate H• Strong acids donate practically all their H’s

HCl + H2O H3O+1 + Cl-1

• Strong bases will react completely with water to form hydroxidesCO3

-2 + H2O HCO3-1 + OH-1

• Weak acids donate a small fraction of their H’s– The process is reversible, the conjugate acid and conjugate base can react

to form the original acid and base

HC2H3O2 + H2O H3O+1 + C2H3O2-1

• Only small fraction of weak base molecules pull H off waterHCO3

-1 + H2O H2CO3 + OH-1

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Figure 15.1: Graphical representation of the behavior of acids in aqueous solution

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Figure 15.2: The relationship of acid strength and conjugate base strength

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Multiprotic Acids

• Monoprotic acids have 1 acid H, diprotic 2, etc.– In oxyacids only the H on the O is

acidic• In strong multiprotic acids, like

H2SO4, only the first H is strong; transferring the second H is usually weak

H2SO4 + H2O H3O+1 + HSO4-1

HSO4-1 + H2O H3O+1 + SO4

-2

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Water as an Acid and a Base

• Amphoteric substances can act as either an acid or a base

– Water as an acid, NH3 + H2O NH4+1 + OH-1

– Water as a base, HCl + H2O H3O+1 + Cl-1

• Water can even react with itself

H2O + H2O H3O +1 + OH-1

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Autoionization of Water• Water is an extremely weak electrolyte

– therefore there must be a few ions present

H2O + H2O H3O+1 + OH-1

• all water solutions contain both H3O+1 and OH-1

– the concentration of H3O+1 and OH-1 are equal

– [H3O+1] = [OH-1] = 10-7M @ 25°C

• Kw = [H3O+1] x [OH-1] = 1 x 10-14 @ 25°C

– Kw is called the ion product constant for water

– as [H3O+1] increases, [OH-] decreases

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Acidic and Basic Solutions

• acidic solutions have a larger [H+1] than [OH-1]

• basic solutions have a larger [OH-1] than [H+1]

• neutral solutions have [H+1]=[OH-1]= 1 x 10-7 M

[H+1] = 1 x 10-14

[OH-1][OH-1] = 1 x 10-14

[H+1]

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Example #2

Determine the given information and the information you need to find

Given [H+1] = 10.0 M Find [OH-1]

Solve the Equation for the Unknown Amount

][H

K ]OH[

][OHx ][H K

1w1-

1-1w

Determine the [H+1] and [OH-1] in a 10.0 M H+1 solution

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Convert all the information to Scientific Notation and Plug the given information into the equation.

Given [H+1] = 10.0 M = 1.00 x 101 M

Kw = 1.0 x 10-14

M 10x 1.0 10x 1.00

10x 1.0 ]OH[

][H

K ]OH[

15-1

14-1-

1w1-

Example #2Determine the [H+1] and [OH-1] in a

10.0 M H+1 solution

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pH & pOH• The acidity/basicity of a solution is often expressed as pH or

pOH• pH = -log[H3O+1] pOH = -log[OH-1]

– pHwater = -log[10-7] = 7 = pOHwater

• [H+1] = 10-pH [OH-1] = 10-pOH

• pH < 7 is acidic; pH > 7 is basic, pH = 7 is neutral• The lower the pH, the more acidic the solution; The higher the pH,

the more basic the solution• 1 pH unit corresponds to a factor of 10 difference in acidity • pOH = 14 - pH

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Figure 15.3: The pH scale and pH values of some common substances

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Figure 15.4: A pH meter

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Figure 15.5: Indicator paper being used to measure the pH of a solution

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Example #3

Find the concentration of [H+1]

M 10x 1.0 10x 1.0

10x 1.0 ]H[

][OH

K ]H[

8-6-

14-1

1w1

Calculate the pH of a solution with a [OH-1] = 1.0 x 10-6 M

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Enter the [H+1] concentration into your calculator and press the log key

log(1.0 x 10-8) = -8.0

Change the sign to get the pHpH = -(-8.0) = 8.0

Example #3

Calculate the pH of a solution with a [OH-1] = 1.0 x 10-6 M

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Enter the [H+1] or [OH-1]concentration into your calculator and press the log key

log(1.0 x 10-3) = -3.0Change the sign to get the pH or pOH

pOH = -(-3) = 3.0Subtract the calculated pH or pOH from

14.00 to get the other valuepH = 14.00 – 3.0 = 11.0

Calculate the pH and pOH of a solution with a [OH-1] = 1.0 x 10-3 M

Example #4

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If you want to calculate [OH-1] use pOH, if you want [H+1] use pH. It may be necessary to convert one to the other using 14 = pH + pOH

pOH = 14.00 – 7.41 = 6.59Enter the pH or pOH concentration into your

calculatorChange the sign of the pH or pOH

-pOH = -(6.59)Press the button(s) on you calculator to take the

inverse log or 10x

[OH-1] = 10-6.59 = 2.6 x 10-7

Example #5Calculate the [OH-1] of a solution with a pH of 7.41

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Calculating the pH of a Strong, Monoprotic Acid

• A strong acid will dissociate 100%

HA H+1 + A-1

• Therefore the molarity of H+1 ions will be the same as the molarity of the acid

• Once the H+1 molarity is determined, the pH can be determined

pH = -log[H+1]

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Example #6

Determine the [H+1] from the acid concentrationHNO3 H+1 + NO3

-1

0.10 M HNO3 = 0.10 M H+1

Enter the [H+1] concentration into your calculator and press the log key

log(0.10) = -1.00Change the sign to get the pH

pH = -(-1.00) = 1.00

Calculate the pH of a 0.10 M HNO3 solution

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Buffered Solutions• Buffered Solutions resist change in pH when an acid

or base is added to it.• Used when need to maintain a certain pH in the

system– Blood

• A buffer solution contains a weak acid and its conjugate base

• Buffers work by reacting with added H+1 or OH-1 ions so they do not accumulate and change the pH

• Buffers will only work as long as there is sufficient weak acid and conjugate base molecules present